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Chapter 8
Periodic
Properties of the
Elements
Electron Configurations
Electron Configuration
of Elements
examples
*end 10/24/16 Lecture
*start 10/26 lecture
after this slide
Valence Electrons
• The electrons in all the sublevels with the highest
principal energy shell are called the valence
electrons.
• Electrons in lower energy shells are called core
electrons.
• One of the most important factors in the way an
atom behaves, both chemically and physically, is
the number of valence electrons.
Electron Configuration
of Valence Elements
examples
Electron Configuration and the
Periodic Table
• The group number corresponds to the
number of valence electrons.
• The length of each “block” is the maximum
number of electrons the sublevel can hold.
• The period number corresponds to the
principal energy level (n) of the valence
electrons.
Irregular Electron Configurations
• We know that, because of sublevel splitting, the 4s sublevel
is lower in energy than the 3d; therefore, the 4s fills before
the 3d.
• But the difference in energy is not large.
• Some of the transition metals have irregular electron
configurations in which the ns only partially fills before the
(n−1)d or doesn’t fill at all.
• Therefore, their electron configuration must be found
experimentally.
Irregular Electron Configurations
•
•
•
•
•
•
Expected
Cr = [Ar]4s23d4
Cu = [Ar]4s23d9
Mo = [Kr]5s24d4
Ru = [Kr]5s24d6
Pd = [Kr]5s24d8
•
•
•
•
•
•
Found experimentally
Cr = [Ar]4s13d5
Cu = [Ar]4s13d10
Mo = [Kr]5s14d5
Ru = [Kr]5s14d7
Pd = [Kr]5s04d10
Properties and Electron Configuration
• The properties of the
elements follow a periodic
pattern.
– Elements in the same column
(same group) have similar
properties.
– The elements in a period show a
pattern that repeats.
• The quantum-mechanical
model explains this because
the number of valence
electrons and the types of
orbitals they occupy are also
periodic.
The Noble Gas Electron Configuration
• The noble gases have eight valence
electrons.
– Except for He, which has only two
electrons
• They are especially nonreactive.
– He and Ne are practically inert.
• The reason the noble gases are so
nonreactive is that the electron
configuration of the noble gases is
especially stable.
The Alkali Metals
• The alkali metals have one more
electron than the previous noble gas.
• In their reactions, the alkali metals
tend to lose one electron, resulting in
the same electron configuration as a
noble gas.
– Forming a cation with a 1+ charge
The Halogens
• Have one fewer electron than the next
noble gas
• In their reactions with metals, the halogens
tend to gain an electron and attain the
electron configuration of the next noble gas,
forming an anion with charge 1−.
• In their reactions with nonmetals, they tend
to share electrons with the other nonmetal
so that each attains the electron
configuration of a noble gas.
Electron Configuration and Ion Charge
• We have seen that many metals and nonmetals form
one ion and that the charge on that ion is predictable
based on its position on the periodic table.
– Group 1A = 1+, group 2A = 2+, group 7A = 1−,
group 6A = 2−, etc.
• These atoms form ions that will result in an electron
configuration that is the same as the nearest noble gas.
Electron Configuration of Anions in Their
Ground State
• Anions are formed when nonmetal atoms gain
enough electrons to have eight valence electrons.
– Filling the s and p sublevels of the valence shell
• The sulfur atom has six valence electrons.
S atom = 1s22s22p63s23p4
• To have eight valence electrons, sulfur must gain
two more.
S2− anion = 1s22s22p63s23p6
Electron Configuration of Cations in Their
Ground State
• Cations are formed when a metal atom loses all its
valence electrons, resulting in a new lower energy level
valence shell.
– However, the process is always endothermic.
• The magnesium atom has two valence electrons.
Mg atom = 1s22s22p63s2
• When magnesium forms a cation, it loses its valence
electrons.
Mg2+ cation = 1s22s22p6
Trend in Atomic Radius: Main Group
• There are several methods for measuring the radius of an
atom, and they give slightly different numbers.
– Van der Waals radius = nonbonding
– Covalent radius = bonding radius
– Atomic radius is an average radius of an atom based on measuring
large numbers of elements and compounds.
Trend in Atomic Radius: Main Group
• Atomic radius decreases across period (left to right).
– Adding electrons to same valence shell
– Effective nuclear charge increases
– Valence shell held closer
Trend in Atomic Radius: Main Group
• Atomic radius (size) increases down group.
– Valence shell farther from nucleus
– Effective nuclear charge fairly close
Periodic Trends in Atomic Radius
Shielding
• In a multielectron system, electrons are simultaneously
attracted to the nucleus and repelled by each other.
• Outer electrons are shielded from the nucleus by the
core electrons.
– Screening or shielding effect
– Outer electrons do not effectively screen for each other.
• The shielding causes the outer electrons to not
experience the full strength of the nuclear charge.
Screening and Effective Nuclear Charge
Electron Configurations of Main Group
Cations in Their Ground State
• Cations form when the atom loses electrons from
the valence shell.
Al atom = 1s22s22p63s23p1
Al3+ ion = 1s22s22p6
Electron Configurations of Transition Metal
Cations in Their Ground State
• When transition metals form cations, the first
electrons removed are the valence electrons, even
though other electrons were added after.
• Electrons may also be removed from the sublevel
closest to the valence shell after the valence
electrons.
• The iron atom has two valence electrons:
Fe atom = 1s22s22p63s23p64s23d6
• When iron forms a cation, it first loses its valence
electrons:
Fe2+ cation = 1s22s22p63s23p63d6
• It can then lose 3d electrons:
Fe3+ cation = 1s22s22p63s23p63d5
Trends in Ionic Radius
• Ions in the same group have the same charge.
• Ion size increases down the column.
– Higher valence shell, larger ion
• Cations are smaller than neutral atoms; anions are
•
larger than neutral atoms.
Cations are smaller than anions.
– Except Rb+ and Cs+, bigger or same size as F− and O2−
• Larger positive charge = smaller cation
– For isoelectronic species
– Isoelectronic = same electron configuration
• Larger negative charge = larger anion
– For isoelectronic species
Periodic Trends in Ionic Radius
Ionization Energy (IE)
• Minimum energy needed to remove an
electron from an atom or ion
–
–
–
–
–
Gas state
Endothermic process
Valence electron easiest to remove, lowest IE
M(g) + IE1 → M1+(g) + 1 e–
M+1(g) + IE2 → M2+(g) + 1 e–
• First ionization energy = energy to remove electron from
neutral atom
• Second IE = energy to remove electron from 1+ ion, etc.
Trends in First Ionization Energy
(larger atom, smaller IE)
• The larger the effective nuclear charge on the
electron, the more energy it takes to remove it.
• The farther the most probable distance the electron
is from the nucleus, the less energy it takes to
remove it.
• First IE decreases down the group.
– Valence electron farther from nucleus
• First IE generally increases across the period.
– Effective nuclear charge increases
Electron Affinity (irregular trends – not responsible for
trends in EA but responsible for definition of EA)
• Energy is released when a neutral atom gains
an electron.
– Gas state
– M(g) + 1e− → M1−(g) + EA
• Electron affinity is defined as exothermic (−)
but may actually be endothermic (+).
– Some alkali earth metals and all noble gases are
endothermic. Why?
• The more energy that is released, the larger the
electron affinity.
– The more negative the number, the larger the EA.
Trends in Electron Affinity
(generally smaller element, larger negative charge EA)
• Alkali metal EA decreases (smaller negative EA)
electron affinity down the column.
– But not all groups do
– Generally irregular increase in EA from second period
to third period
• “Generally” increases (larger negative EA) across
period
– Becomes more negative from left to right
– Not absolute
– Group 5A generally lower EA than expected because
extra electron must pair
– Groups 2A and 8A generally very low EA because added
electron goes into higher energy level or sublevel
• Highest EA in any period = halogen
Properties of Metals and Nonmetals
• Metals
–
–
–
–
–
–
Malleable and ductile
Shiny, lustrous, reflect light
Conduct heat and electricity
Most oxides basic and ionic
Form cations in solution
Lose electrons in reactions—oxidized
• Nonmetals
–
–
–
–
–
–
Brittle in solid state
Dull, nonreflective solid surface
Electrical and thermal insulators
Most oxides are acidic and molecular
Form anions and polyatomic anions
Gain electrons in reactions—reduced
Metallic Character
• Metallic character is how closely an element’s
properties match the ideal properties of a metal.
– More malleable and ductile, better conductors, and easier
to ionize
• Metallic character decreases left to right across a
period.
– Metals found at the left of the period, and nonmetals to
the right
• Metallic character increases down the column.
– Nonmetals found at the top of the middle main-group
elements, and metals found at the bottom
End 10/26/16
Alkali Metals (more reactive down group)
Trends in the Halogens (more reactive down group)
• Reactivity increases down the column.
• They react with hydrogen to form HX, acids.
• Melting point and boiling point increase down the
column.
• Density increases down the column.
– In general, the increase in mass is greater than the
increase in volume.
Supplementary Material:
Quantum-Mechanical Explanation for the
Group Trend in Atomic Radius
• The size of an atom is related to the distance the
valence electrons are from the nucleus.
• The larger the orbital an electron is in, the farther
its most probable distance will be from the nucleus
and the less attraction it will have for the nucleus.
Quantum-Mechanical Explanation for the
Group Trend in Atomic Radius
• Traversing down a group adds a principal energy level.
• The larger the principal energy level an orbital is in, the
larger its volume.
• Quantum-mechanics predicts the atoms should get
larger down a column.
Quantum-Mechanical Explanation for the
Period Trend in Atomic Radius
• The larger the effective nuclear charge an electron
experiences, the stronger the attraction it will have
for the nucleus.
• The stronger the attraction the valence electrons
have for the nucleus, the closer their average
distance will be to the nucleus.
• Traversing across a period increases the effective
nuclear charge on the valence electrons.
• Quantum-mechanics predicts the atoms should get
smaller across a period.
Trends in Atomic Radius: Transition Metals
• Atoms in the same group increase in size down the
column.
• Atomic radii of transition metals are roughly the
same size across the d block.
– Much less difference than across main-group elements
– Valence shell ns2, not the (n−1)d electrons
– Effective nuclear charge on the ns2 electrons
approximately the same
Magnetic Properties of Transition Metal
Atoms and Ions
• Electron configurations that result in unpaired
electrons mean that the atom or ion will have a net
magnetic field; this is called paramagnetism.
– Will be attracted to a magnetic field
• Electron configurations that result in all paired
electrons mean that the atom or ion will have no
magnetic field; this is called diamagnetism.
– Slightly repelled by a magnetic field
Explanation for the Trends in Cation Radius
• When atoms form cations, the valence electrons are
removed.
• The farthest electrons from the nucleus are the p or d
electrons in the (n − 1) energy level.
• This results in the cation being smaller than the atom.
Explanation for the Trends in Cation Radius
• These “new valence electrons” also experience a
larger effective nuclear charge than the “old
valence electrons,” shrinking the ion even more.
• Traversing down a group increases the (n − 1)
level, causing the cations to get larger.
• Traversing to the right across a period increases
the effective nuclear charge for isoelectronic
cations, causing the cations to get smaller.
Periodic Trends in Anionic Radius
Explanation for the Trends in Anion Radius
• When atoms form anions, electrons are added to
the valence shell.
• These “new valence electrons” experience a
smaller effective nuclear charge than the “old
valence electrons,” increasing the size.
• The result is that the anion is larger than the atom.
Explanation for the Trends in Anion Radius
• Traversing down a group increases the n level,
causing the anions to get larger.
• Traversing to the right across a period decreases
the effective nuclear charge for isoelectronic
anions, causing the anions to get larger.
Explanation for the Trends in Metallic
Character
• Metals generally have smaller first ionization
energies, and nonmetals generally have larger
electron affinities.
– Except for the noble gases
• ∴ quantum mechanics predicts the atom’s metallic
character should increase down a column because
the valence electrons are not held as strongly.
• ∴ quantum mechanics predicts the atom’s metallic
character should decrease across a period
because the valence electrons are held more
strongly and the electron affinity increases.
Trends in the Alkali Metals
• Atomic radius increases down the column.
• Ionization energy decreases down the column.
• Very low ionization energies
–
–
–
–
Good reducing agents; easy to oxidize
Very reactive; not found uncombined in nature
React with nonmetals to form salts
Compounds generally soluble in water ∴ found in
seawater
• Electron affinity decreases down the column.
• Melting point decreases down the column.
– All very low MP for metals
• Density increases down the column.
– Except K
– In general, the increase in mass is greater than the
increase in volume.
Trends in the Halogens
• Atomic radius increases down the column.
• Ionization energy decreases down the column.
• Very high electron affinities
–
–
–
–
Good oxidizing agents; easy to reduce
Very reactive; not found uncombined in nature
React with metals to form salts
Compounds generally soluble in water ∴ found in
seawater
Halogens
Reactions of Alkali Metals with Halogens
• Alkali metals are oxidized
to the 1+ ion.
• Halogens are reduced to
the 1− ion.
• The ions then attach
together by ionic bonds.
• The reaction is
exothermic.
Reactions of Alkali Metals with Water
• Alkali metals are oxidized to the 1+ ion.
• H2O is split into H2(g) and OH− ion.
• The Li, Na, and K are less dense than the water, so they
float on top.
• The ions then attach together by ionic bonds.
• The reaction is exothermic, and often the heat released
ignites the H2(g).
Trends in the Noble Gases
• Atomic radius increases down the column.
• Ionization energy decreases down the column.
– Very high IE
• Very unreactive
– Only found uncombined in nature
– Used as “inert” atmosphere when reactions with other
gases would be undesirable
Trends in the Noble Gases
• Melting point and boiling point increase down the
column.
– All gases at room temperature
– Very low boiling points
• Density increases down the column.
– In general, the increase in mass is greater than the
increase in volume.
Noble Gases
Exceptions in the First IE Trends
• First ionization energy generally increases from
•
left to right across a period.
Except from 2A to 3A and 5A to 6A
Exceptions in the First Ionization Energy
Trends, N and O
To ionize N, you must break up a half-full sublevel, which costs
extra energy.
When you ionize O, you get a half-full sublevel, which costs
less energy.
Trends in Successive Ionization Energies
• Removal of each successive
electron costs more energy.
– Shrinkage in size due to having
more protons than electrons
– Outer electrons closer to
the nucleus; therefore
harder to remove
• There’s a regular increase in
energy for each successive
valence electron.
• There’s a large increase in
energy when core electrons
are removed.
Trends in Second and Successive Ionization
Energies